Chapter 1 Chemistry: Matter on the Atomic Scale

Transcription

1 Chapter 1 In This Chapter This chapter introduces the fundamental components of matter, the different types of structures they can make when they join together, and the types of changes they undergo. Chapter Outline 1.1 What Is Chemistry? 1.2 Classification of Matter 1.3 Units and Measurement 1.4 Unit Conversions Chapter Review Chapter Summary Assignment 1.1 What Is Chemistry? Section Outline 1.1a The Scale of Chemistry 1.1b Measuring Matter Section Summary Assignment This section introduces you to chemistry, the study of matter, its transformations, and how it behaves. We define matter as any physical substance that occupies space and has mass. Matter consists of atoms and molecules, and it is at the atomic and molecular level that chemical transformations take place. Opening Exploration 1.1: Levels of Matter 1.1a The Scale of Chemistry Different fields of science examine the world at different levels of detail (Interactive Figure 1.1.1). 1

2 Interactive Figure Understand the scale of science. The macroscopic, microscopic, and atomic scales in different fields of science. When describing matter that can be seen with the naked eye, scientists are working on the macroscopic scale. Chemists use the atomic scale (sometimes called the nanoscale or the molecular scale) when describing individual atoms or molecules. In general, in chemistry we make observations at the macroscopic level and we describe and explain chemical processes on the atomic level. That is, we use our macroscopic scale observations to explain atomic scale properties. 1.1b Measuring Matter Chemistry is an experimental science that involves designing thoughtful experiments and making careful observations of macroscopic amounts of matter. Everything that is known about how atoms and molecules interact has been learned through making careful observations on the macroscopic scale and inferring what those observations must mean about atomic scale objects. For example, careful measurement of the mass of a chemical sample before and after it is heated provides information about the chemical composition of a substance. Observing how a chemical sample behaves in the presence of a strong magnetic field such as that found in a magnetic resonance imaging (MRI) scanner provides information about how molecules and atoms are arranged in human tissues. An important part of chemistry and science in general is the concept that all ideas are open to challenge. When we perform measurements on chemical substances and interpret the results in terms of atomic scale properties, the results are always examined to see if there are alternative ways to interpret the data. This method of investigation leads to chemical information about the properties and behavior of matter that is supported by the results of many different experiments

3 EXAMPLE PROBLEM: Differentiate between the macroscopic and atomic scales. Classify each of the following as matter that can be measured or observed on either the macroscopic or atomic scale. (a) An RNA molecule (b) A mercury atom (c) A sample of liquid mercury SOLUTION: You are asked to identify whether a substance can be measured or observed on the macroscopic or atomic scales. You are given the identity of the substance. (a) Atomic scale. An RNA molecule is too small to be seen with the naked eye or with an optical microscope. (b) Atomic scale. Individual atoms cannot be seen with the naked eye or with an optical microscope. (c) Macroscopic scale. Liquid mercury can be seen with the naked eye T: Tutorial Assignment Differentiate Between the Macroscopic and Atomic Scales: Tutorial Assignment 1.1.1: Mastery Assignment Differentiate Between the Macroscopic and Atomic Scales: Mastery Assignment 1.2 Classification of Matter Section Outline 1.2a Classifying Matter on the Atomic Scale 1.2b Classifying Pure Substances on the Macroscopic Scale 1.2c Classifying Mixtures on the Macroscopic Scale Section Summary Assignment Matter can be described by a collection of characteristics called properties. One of the fundamental properties of matter is its composition, or the specific types of atoms or molecules that make it up. In this section, we describe investigate the properties of pure substances such as elements and compounds and of mixtures, which contain more than one substance

4 Opening Exploration 1.2 Separating Mixtures 1.2a Classifying Matter on the Atomic Scale An element, which is the simplest type of matter, is a pure substance that cannot be broken down or separated into simpler substances. [Flashforward 1.1 anchor] You are already familiar with some of the most common elements such as gold, silver, and copper, which are used in making coins and jewelry, and oxygen, nitrogen, and argon, which are the three most abundant gases in our atmosphere. A total of 116 elements have been identified, 90 of which exist in nature (the rest have been synthesized in the laboratory). Elements are represented by a one- or two-letter element symbol, and they are organized in the periodic table that is shown in Chapter 2 and in the Reference Tools. A few common elements and their symbols are shown in Table Notice that the when the symbol for an element consists of two letters, only the first letter is capitalized. Flashforward 1.1 Information about the elements is organized in the periodic table of the elements the most important tool that chemists use. Not only does this table contain information specific to each element, but it also organizes the elements according to their physical and chemical properties. (See Section 2.2.) Table Some Common Elements and Their Symbols Name Hydrogen Carbon Oxygen Sodium Iron Aluminum Symbol H C O Na Fe Al - 4 -

5 Atoms An atom is the smallest indivisible unit of an element. For example, the element aluminum is made up entirely of aluminum atoms, as shown in Interactive Figure Interactive Figure Explore the composition of a sample of aluminum. A piece of aluminum and a representation of aluminum at the molecular level. Although individual atoms are too small to be seen directly with the naked eye or with the use of a standard microscope, methods such as scanning tunneling microscopy (STM) allow scientists to view atoms. Both experimental observations and theoretical studies show that isolated atoms are spherical and that atoms of different elements have different sizes. Thus, the model used to represent isolated atoms consists of spheres of different sizes. In addition, chemists often use color to distinguish atoms of different elements. For example, oxygen atoms are usually represented as red spheres; carbon atoms, as gray or black spheres; and hydrogen atoms, as white spheres. Elements are made up of only one type of atom. For example, the element oxygen is found in two forms: as O2, in which two oxygen atoms are grouped together, and as O3, in which three oxygen atoms are grouped together. The most common form of oxygen is O2, dioxygen, a gas that makes up about 22% of the air we breathe. Ozone, O3, is a gas with a distinct odor that can be toxic to humans. Both dioxygen and ozone are elemental forms of oxygen because they consist of only one type of atom

6 Compounds and Molecules A chemical compound is a substance formed when two or more elements are combined in a defined ratio. Compounds differ from elements in that they can be broken down chemically into simpler substances. You have encountered chemical compounds in many common substances, such as table salt, a compound consisting of the elements sodium and chlorine, and phosphoric acid, a compound found in soft drinks that contain hydrogen, oxygen, and phosphorus. Molecules are collections of atoms that are held together by chemical bonds. In models used to represent molecules, chemical bonds are often represented using cylinders or lines that connect atoms, represented as spheres. The composition and arrangement of elements in molecules affects the properties of a substance. For example, molecules of both water (H 2 O) and hydrogen peroxide ( H2O2 ) contain only the elements hydrogen and oxygen. Water is a relatively inert substance that is safe to drink in its pure form. Hydrogen peroxide, however, is a reactive liquid that is used to disinfect wounds and can cause severe burns if swallowed. EXAMPLE PROBLEM: Classify pure substances as elements or compounds. Classify each of the following substances as either an element or a compound. (a) Si (b) CO 2 (c) P 4 SOLUTION: You are asked to classify a substance as an element or a compound. You are given the chemical formula of the substance. (a) Element. Silicon is an example of an element because it consists of only one type of atom. (b) Compound. This compound contains both carbon and oxygen. (c) Element. Although this is an example of a molecular substance, it consists of only a single type of atom T: Tutorial AssignmentClassify Pure Substances as Elements or Compounds: Tutorial Assignment 1.2.1: Mastery Assignment Classify Pure Substances as Elements or Compounds: Mastery Assignment - 6 -

7 1.2b Classifying Pure Substances on the Macroscopic Scale A pure substance contains only one type of element or compound and has fixed chemical composition. A pure substance also has characteristic properties, measurable qualities that are independent of the sample size. The physical properties of a chemical substance are those that do not change the chemical composition of the material when they are measured. Some examples of physical properties include physical state, color, viscosity (resistance to flow), opacity, density, conductivity, and melting and boiling points. States of Matter One of the most important physical properties is the physical state of a material. The three physical states of matter are solid, liquid, and gas (Interactive Figure 1.2.2). Interactive Figure Distinguish the properties of the three states of matter. Representations of a solid, a liquid, and a gas. The macroscopic properties of these states are directly related to the arrangement and properties of particles at the atomic level. At the macroscopic level, a solid is a dense material with a defined shape. At the atomic level, the atoms or molecules of a solid are packed together closely. The atoms or molecules are vibrating, but they do not move past one another. At the macroscopic level, a liquid is also dense, but unlike a solid it flows and takes on the shape of its container. At the atomic level, the atoms or molecules of a liquid are close together, but they move more than the particles in a solid and can flow past one another. Finally, at the macroscopic level, a gas has no fixed shape or volume. At the atomic level, the atoms or molecules of a gas are spaced widely apart and are moving rapidly past one another. The particles of a gas do not strongly interact with one another, and they move freely until they collide with one another or with the walls of the container. The physical state of a substance can change when energy, often in the form of heat, is added or removed. When energy is added to a solid, the temperature at which the solid is converted to a liquid is the melting point of the substance. The conversion of liquid to solid occurs at the same temperature as energy is removed (the temperature falls) and is called the freezing point. A liquid is converted to a gas at the boiling point of a substance. As you will see in the following section, melting and boiling points are measured in Celsius (ºC) or Kelvin (K) temperature units. Not all materials can exist in all three physical states. Polyethylene, for example, does not exist as a gas. Heating a solid polyethylene milk bottle at high temperatures causes it to decompose into other substances. Helium, a gas at room temperature, can be liquefied at very low temperatures, but it is not possible to solidify helium. A change in the physical property of a substance is called a physical change. Physical changes may change the appearance or the physical state of a substance, but they do not change its chemical - 7 -

8 composition. For example, a change in the physical state of water changing from a liquid to a gas involves a change in how the particles are packed together at the atomic level, but it does not change the chemical makeup of the material. Chemical Properties The chemical properties of a substance are those that involve a chemical change in the material and often involve a substance interacting with other chemicals. For example, a chemical property of methanol, CH 3 OH, is that it is highly flammable because the compound burns in air (it reacts with oxygen in the air) to form water and carbon dioxide (Interactive Figure 1.2.3). Interactive Figure Investigate the chemical properties of methanol. [Figure ID #1-9] Methanol is a flammable liquid. A chemical change involves a change in the chemical composition of the material. The flammability of methanol is a chemical property, and demonstrating this chemical property involves a chemical change. EXAMPLE PROBLEM Identify physical and chemical properties and physical and chemical changes. (a) When aluminum foil is placed into liquid bromine a white solid forms. Is this a chemical or physical property of aluminum? (b) Iodine is a purple solid. Is this a chemical or physical property of iodine? (c) Classify each of the following changes as chemical or physical. Boiling water Baking bread SOLUTION: You are asked to identify a change or property as chemical or physical. You are given a description of a material or a change. (a) Chemical property. Chemical properties are those that involve a chemical change in the material and often involve a substance interacting with other chemicals. In this example, one substance (the aluminum) is converted into a new substance (a white solid). (b) Physical property. A physical property such as color is observed without changing the chemical identity of the substance. (c) (i) Physical change.a physical change alters the physical form of a substance without changing its chemical identity. Boiling does not change the chemical composition of water. (ii) Chemical change. When a chemical change takes place, the original substances (the bread ingredients) are broken down and a new substance (bread) is formed T: Tutorial Assignment Identify Physical and Chemical Properties and Physical and Chemical Changes: Tutorial Assignment 1.2.2: Mastery Assignment Identify Physical and Chemical Properties and Physical and Chemical Changes: Mastery Assignment - 8 -

9 1.2c Classifying Mixtures on the Macroscopic Scale As you can see when you look around you, the world is made of complex materials. Much of what surrounds us is made up of mixtures of different substances. A mixture is a substance made up of two or more elements or compounds that have not reacted chemically. Unlike compounds, where the ratio of elements is fixed, the relative amounts of different components in a mixture can vary. Mixtures that have a constant composition throughout the material are called homogeneous mixtures. For example, dissolving table salt in water creates a mixture of the two chemical compounds water (H 2 O) and table salt (NaCl). Because the mixture is uniform, meaning that the same ratio of water to table salt is found no matter where it is sampled, it is a homogeneous mixture. A mixture in which the composition is not uniform is called a heterogeneous mixture. For example, a cold glass of freshly squeezed lemonade with ice is a heterogeneous mixture because you can see the individual components (ice cubes, lemonade, and pulp) and the relative amounts of each component will depend on where the lemonade is sampled (from the top of the glass or from the bottom). The two different types of mixtures are explored in Interactive Figure Interactive Figure Identify homogeneous and heterogeneous mixtures. Homogeneous and heterogeneous mixtures. Homogeneous and heterogeneous mixtures can usually be physically separated into individual components. For example, a homogeneous mixture of salt and water is separated by heating the mixture to evaporate the water, leaving behind the salt. A heterogeneous mixture of sand and water is separated by pouring the mixture through filter paper. The sand is trapped in the filter while the water passes through. Heating the wet sand to evaporate the remaining water completes the physical separation. Like pure substances, mixtures have physical and chemical properties. These properties, however, depend on the composition of the mixture. For example, a mixture of 10 grams of table sugar and 100 grams of water has a boiling point of ºC while a mixture of 20 grams of table sugar and 100 grams of water has a boiling point of ºC. Interactive Figure summarizes how we classify different forms of matter in chemistry

10 Interactive Figure Classify matter. A flow chart for the classification of matter EXAMPLE PROBLEM: Identify pure substances and mixtures. Classify each of the following as a pure substance, a homogeneous mixture, or a heterogeneous mixture. (a) Copper wire (b) Oil and vinegar salad dressing (c) Vinegar SOLUTION: You are asked to classify items as a pure substance, a heterogeneous mixture, or a homogeneous mixture. You are given the identity of the item. (a) Pure substance. Copper is an element. (b) Heterogeneous mixture. The salad dressing is a mixture that does not have a uniform composition. The different components are visible to the naked eye, and the composition of the mixture varies with the sampling location. (c) Homogeneous mixture. Vinegar is a uniform mixture of water, acetic acid, and other compounds. The different components in this mixture are not visible to the naked eye T: Tutorial Assignment Identify Pure Substances and Mixtures: Tutorial Assignment 1.2.3: Mastery Assignment Identify Pure Substances and Mixtures: Mastery Assignment

11 1.3 Units and Measurement Section Outline 1.3a Scientific Units 1.3b Scientific Notation 1.3c SI Base Units: Length, Mass, and Temperature 1.3d SI-Derived Units: Volume, Density, and Energy 1.3e Significant Figures 1.3f Significant Figures and Calculations 1.3g Precision and Accuracy Section Summary Assignment Chemistry involves observing matter, and our observations are substantiated by careful measurements of physical quantities. Chemists in particular need to make careful measurements because we use those measurements to infer the properties of matter on the atomic scale. In this section, we explore how chemists work with the numbers and units associated with these measurements. Opening Exploration 1.3 Scientific Units and Everyday Objects [Figure ID#1-12] 1.3a Scientific Units Some of the most common measurements in chemistry are mass, volume, time, temperature, and density. Measuring these quantities allows us to describe the chemical and physical properties of matter and study the chemical and physical changes that matter undergoes. When reporting a measurement, we use scientific units to indicate what was measured. SI units, abbreviated from the French Système International d Unités, are used in scientific measurements in almost all countries. This unit system consists of seven base units. Other units are called derived units and are combinations of the base units (Table 1.3.1) Table SI Units Physical Quantity SI Unit Symbol Base Units Length Meter m Mass Kilogram kg Time Second s Electric current Ampere A Temperature Kelvin K Amount of substance Mole mol Luminous intensity Candela cd Some Derived Units Volume m 3 (cubic meter) Density kg/m 3 Energy J (joule)

12 Metric prefixes are combined with SI units when reporting physical quantities in order to reflect the relative size of the measured quantity. Table shows the metric prefixes most commonly used in scientific measurements. Table Common Prefixes Used in the SI and Metric Systems Prefix Abbreviation Factor mega M 100,000 (10 6 ) kilo k 1000 (10 3 ) deci d 0.1 (10 1 ) centi c 0.01 (10 2 ) milli m (10 3 ) micro (10 6 ) nano n (10 9 ) pico p (10 12 ) When making and reporting measurements, it is important to use both the value and the appropriate units. For example, in the United States, speed limits are reported in units of miles per hour (mph). A U.S. citizen traveling to Canada might see a speed limit sign reading 100 and assume the units are mph. This could be an expensive mistake, however, because speed limits in Canada are reported in units of kilometers per hour (km/h); a 100 km/h speed limit is the equivalent of 62 mph. Quick Check What is the SI unit for mass? a. gram (g) b. kilogram (kg) c. pound (lb) 2. Which is longer, 1 mm or 1 nm? a. 1 mm b. 1 nm 1.3b Scientific Notation Numbers that are very large or very small can be represented using scientific notation. A number written in scientific notation has the general form N 10 x, where N is a number between 1 and 10 and x is a positive or negative integer. For example, the number is written as and the number is written as in scientific notation. Notice that x is positive for numbers greater than 1 and negative for numbers less than 1. To convert a number from standard notation to scientific notation, count the number of times the decimal point must be moved to the right (for numbers less than 1) or to the left (for numbers greater than 1) in order to result in a number between 1 and 10. For the number 13433,

13 the decimal point is moved four places to the left and the number is written When a number is less than 1, the decimal point is moved to the right and the exponent (x) is negative. For the number , the decimal point is moved three places to the right and the number is written Notice that in both cases, moving the decimal point one place is the equivalent of multiplying or dividing by 10. To convert a number from exponential notation to standard notation, write the value of N and then move the decimal point x places to the right if x is positive or move the decimal point x places to the left if x is negative. EXAMPLE PROBLEM:Write numbers using scientific notation. a) Write the following numbers in scientific notation: (i) (ii) (b) Write the following numbers in standard notation: (i) (ii) SOLUTION: You are asked to convert between standard and scientific notation. You are given a number in standard or scientific notation. (a) (i) Moving the decimal point five places to the right results in a number between 1 and 10. The exponent is negative because this number is less than (ii) Moving the decimal point nine places to the left results in a number between 1 and 10. The exponent is positive because this number is greater than (b) (i) Move the decimal point six places to the right (ii) Move the decimal point three places to the left T: Tutorial Assignment Write Numbers Using Scientific Notation: Tutorial Assignment 1.3.1: Mastery Assignment Write Numbers Using Scientific Notation: Mastery Assignment 1.3c SI Base Units: Length, Mass, and Temperature Length The SI unit of length is the meter, m. A pencil has a length of about 0.16 m, which is equivalent to 16 centimeters (cm). Atomic radii can be expressed using nanometer (nm) or picometer (pm) units. The definition of the meter is based on the speed of light in a vacuum, exactly 299,792,458 meters per second. One meter is therefore the length of the path traveled by light in vacuum during 1/299,792,458 of a second. Mass The SI unit of mass is the kilogram, kg. This is the only SI base unit that contains a metric prefix. One kilogram is equal to approximately 2.2 pounds (lb). In the chemistry lab, the mass of a

14 sample is typically measured using units of grams (g) or milligrams (mg). The kilogram standard is the mass of a piece of platinum-iridium alloy that is kept at the International Bureau of Weights and Measures. Temperature Temperature is a relative measure of how hot or cold a substance is and is commonly reported using one of three temperature scales. In the United States, temperatures are commonly reported using the Fahrenheit temperature scale that has units of degrees Fahrenheit (ºF). In scientific measurements, the Celsius and Kelvin temperature scales are used, with units of degrees Celsius (ºC) and kelvins (K), respectively. Notice that for the Kelvin temperature scale, the name of the temperature unit (kelvin) is not capitalized but the abbreviation, K, is capitalized. As shown in Interactive Figure 1.3.1, the three temperature scales have different defined values for the melting and freezing points of water. Interactive Figure Compare different temperature scales. Fahrenheit, Celsius, and Kelvin temperature scales In the Fahrenheit scale, the freezing point of water is set at 32 ºF and the boiling point is 180 degrees higher, 212 ºF. In the Celsius temperature scale, the freezing point of water is assigned a temperature of 0 ºC and the boiling point of water is assigned a temperature of 100 ºC. The lowest temperature on the Kelvin temperature scale, 0 K, is degrees lower than 0 ºC. This temperature, known as absolute zero, is the lowest temperature possible. The Celsius and Kelvin temperature scales are similar in that a 1-degree increment is the same on both scales. That is, an increase of 1 K is equal to an increase of 1 ºC. Equation 1.1 shows the relationship between the Celsius and Kelvin temperature scales. T (ºC) = T (K) (1.1)

15 The Fahrenheit and Celsius temperature scales differ in the size of a degree. 180 Fahrenheit degrees = 100 Celsius degrees 9 5 Fahrenheit degrees = 1 Celsius degree Equation 1.2 shows the relationship between the Fahrenheit and Celsius temperature scales: 9 T (ºF) = 5 [T (ºC)] + 32 (1.2) EXAMPLE PROBLEM Interconvert Fahrenheit, Celsius, and Kelvin temperatures. The boiling point of a liquid is K. What is this temperature on the Celsius and Fahrenheit scales? SOLUTION: You are asked to convert a temperature from kelvin to Celsius and Fahrenheit units. You are given a temperature in kelvin units. Convert the temperature to Celsius temperature units. T (ºC) = T (K) T (ºC) = K = ºC Use the temperature in Celsius units to calculate the temperature on the Fahrenheit scale. 9 T (ºF) = 5 [T (ºC)] T (ºF) = 5 (82.63 ºC) + 32 T (ºF) = ºF Is your answer reasonable? The Celsius temperature should be greater than zero because kelvin temperature is greater than K, which is equal to 0 ºC. The Celsius temperature is close to 100 ºC, the boiling point of water, so it is reasonable for the Fahrenheit temperature to be ºF because this is close to the boiling point of water on the Fahrenheit scale (212 ºF) T: Tutorial Assignment Interconvert Fahrenheit, Celsius, and Kelvin Temperatures: Tutorial Assignment 1.3.2: Mastery Assignment Interconvert Fahrenheit, Celsius, and Kelvin Temperatures: Mastery Assignment 1.3d SI Derived Units: Volume, Density, and Energy Volume Although the SI unit of volume is the cubic meter (m 3 ), a more common unit of volume is the liter (L). Notice that the abbreviation for liter is a capital L. A useful relationship to remember is that one milliliter is equal to one cubic centimeter (1 ml = 1 cm 3 ). Energy The SI unit of energy is the joule (J). Another common energy unit is the calorie (cal), and one calorie is equal to joules (1 cal = J). One dietary calorie (Cal) is equal to 1000 calories (1 Cal = 1000 cal = 1 kcal)

16 Density The density of a substance is a physical property that relates the mass of a substance to its volume (Equation 1.3). mass density = volume (1.3) The densities of solids and liquids are reported in units of grams per cubic centimeter (g/cm 3 ) or grams per milliliter (g/ml), whereas the density of a gas is typically reported in units of grams per liter (g/l). Because the volume of most substances changes with a change in temperature, density also changes with temperature. Most density values are reported at a standard temperature, 25 ºC, close to room temperature. The densities of some common substances are listed in Interactive Table Interactive Table Densities of Some Common Substances at 25 ºC Substance Density (g/cm 3 ) Iron 7.86 Sodium chloride 2.16 Water 1.00 Oxygen Density can be calculated from mass and volume data as shown in the following example. EXAMPLE PROBLEM: CALCULATE DENSITY. A 5.78-mL sample of a colorless liquid has a mass of 4.54 g. Calculate the density of the liquid and identify it as either ethanol (density = g/ml) or benzene (density = g/ml). SOLUTION: You are asked to calculate the density of a liquid and identify the liquid. You are given the mass and volume of a liquid and the density of two liquids. Use Equation 1.3 to calculate the density of the liquid. mass 4.54 g density = volume = 5.78 ml = g/ml The liquid is ethanol T: Tutorial Assigment Calculate Density: Tutorial Assignment 1.3.3: Mastery Assignment Calculate Density: Mastery Assignment 1.3e Significant Figures The certainty in any measurement is limited by the instrument that is used to make the measurement. For example, an orange weighed on a grocery scale weighs 157 g. A standard laboratory balance, however, like the one shown in Figure 1.3.2, reports the mass of the same orange as g [Figure ID#1-16] Figure A top loading laboratory balance

17 In both cases, some uncertainty is present in the measurement. The grocery scale measurement is certain to the nearest 1 g, and the value is reported as 157 ± 1 g. The laboratory scale measurement has less uncertainty, and the mass of the orange is reported as ± g. In general, we will drop the ± symbol and assume an uncertainty of one unit in the rightmost digit when reading a measurement. When using a nondigital measuring device such as a ruler or a graduated cylinder, we always estimate the rightmost digit when reporting the measured value. A digital measuring device such as a top-loading laboratory balance or ph meter includes the estimated digit in its readout. Some measured quantities are infinitely certain, or exact. For example, the number of oranges you have purchased at the grocery store is an exact number. Some units are defined with exact numbers, such as the metric prefixes (1 mm = m) and the relationship between inches and centimeters (1 in = 2.54 cm, exactly). The digits in a measurement, both the certain and uncertain digits, are called significant figures or significant digits. For example, the mass of an orange has three significant figures when measured using a grocery scale (157 g) and six significant figures when measured on a laboratory balance ( g). Some simple rules are used to determine the number of significant figures in a measurement (Interactive Table 1.3.4). For example, consider the numbers and Interactive Table Rules for Determining Significant Figures 1. All non-zero digits and zeros between non-zero digits are significant. In , the digits 3 and 8 and the zero between 3 and 8 are significant. In , the digits 7, 2, 8 and 6 and the zero between 8 and 6 are significant. 2. In numbers containing a decimal point, (a) all zeros at the end of the number are significant. In , the zero to the right of 8 is significant. (b) all zeros at the beginning of the number are not significant. In , both zeros to the left of 3 are not significant. 3. In numbers with no decimal point, all zeros at the end of the number are not significant. In , the zero to the right of 6 is not significant. Thus, the number has four significant figures and the number has five significant figures. Notice that for numbers written in scientific notation, the number of significant figures is equal to the number of digits in the number written before the exponent. For example, the number has three significant figures and has four significant figures

18 . EXAMPLE PROBLEM: Identify the significant figures in a number Identify the number of significant figures in the following numbers. (a) (b) (c) 1030 SOLUTION: You are asked to identify the number of significant figures in a number. You are given a number. (a) All non-zero digits are significant (there are four), and because this number has a decimal point, the zeros at the end of the number are also significant (there are two). This number has six significant figures. (b) All non-zero digits are significant (there are two), and because this number has a decimal point, the three zeros at the beginning of the number are not significant. This number has two significant figures. (c) All non-zero digits are significant (there are two), and the zero between the non-zero digits is also significant (there is one). Because this number has no decimal, the zero at the end of the number is not significant. This number has three significant figures T: Tutorial Assignment Identify the Significant Figures in a Number: Tutorial Assignment 1.3.4: Mastery Assignment Identify the Significant Figures in a Number: Mastery Assignment 1.3f Significant Figures and Calculations When calculations involving measurements are performed, the final calculated result is no more certain than the least certain number in the calculation. If necessary, the answer is rounded to the correct number of significant figures. For example, consider a density calculation involving a sample with a mass of 3.2 g and a volume of cm 3. A standard calculator reports a density of g/cm 3, but is this a reasonable number? In this case, the least certain number is the sample mass with two significant figures and the calculated density therefore has two significant figures, or 0.12 g/cm 3. The final step in a calculation usually involves rounding the answer so that it has the correct certainty. When numbers are rounded, the last digit retained is increased by 1 only if the digit that follows is 5 or greater. When you are performing calculations involving multiple steps, it is best to round only at the final step in the calculation. That is, carry at least one extra significant figure during each step of the calculation to minimize rounding errors in the final calculated result. When multiplication or division is performed, the certainty in the answer is related to the significant figures in the numbers in the calculation. The answer has the same number of significant figures as the measurement with the fewest significant figures. For example, = Round off to 265 Number of significant figures: = Round off to Number of significant figures:

19 When addition or subtraction is performed, the certainty of the answer is related to the decimal places present in the numbers in the calculation. The answer has the same number of decimal places as the number with the fewest decimal places. For example, = Round off to Number of decimal places: = Round off to Number of decimal places: Exact numbers do not limit the number of significant figures or the number of decimal places in a calculated result. For example, suppose you want to convert the amount of time it takes for a chemical reaction to take place, 7.2 minutes, to units of seconds. There are exactly 60 seconds in 1 minute, so it is the number of significant figures in the measured number that determines the number of significant figures in the answer. 7.2 minutes 60 seconds/minute = 430 seconds Number of significant figures: 2 Exact number

20 EXAMPLE PROBLEM: Use significant figures in calculations. Carry out the following calculations, reporting the answer using the correct number of significant figures. (a) (8.3145)( ) (b) SOLUTION: You are asked to carry out a calculation and report the answer using the correct number of significant figures. You are given a mathematical operation. a) The numbers in this multiplication operation have five (8.3145) and two ( ) significant figures. Therefore, the answer should be reported to two significant figures: (8.3145)( ) = (b) This calculation involves both addition and division, so it must be completed in two steps. First, determine the number of significant figures that result from the addition operation in the numerator of this fraction. The first value, 25, has no decimal places, and the second value, , has two decimal places. Therefore, the sum should have no decimal places and, as a result, has three significant figures = 298 Now the division operation can be performed. The numbers in this operation have three (298) and four (1.750) significant figures. Therefore, the answer should be reported to three significant figures = When doing multi-step calculations such as this one, it is a good idea to identify the correct number of significant figures that result from each operation but to round to the correct number of significant figures only in the last step of the calculation. This will help minimize rounding errors in your answers T: Tutorial Assignment Use Significant Figures in Calculations: Tutorial Assignment 1.3.5: Mastery Assignment Use Significant Figures in Calculations: Mastery Assignment 1.3g Precision and Accuracy Along with the number of significant figures in a value, the certainty of a measurement can also be described using the terms precision and accuracy. Precision is how close the values in a set of measurements are to one another. Accuracy is how close a measurement or a set of measurements is to a real value. Interactive Figure demonstrates the relationship between precision and accuracy

21 Interactive Figure Distinguish between precision and accuracy. These darts were thrown with precision but without accuracy. 1.4 Unit Conversions Section Outline 1.4a Dimensional Analysis 1.4b Unit Conversions Using Density Section Summary Assignment Quantities of matter, energy, or time can be expressed using different units. In this section, we describe dimensional analysis, the method by which we can convert the representation of a quantity from one unit (or set of units) to another. Opening Exploration 1.4 How Important Are Units? [Figure ID # 1-18] 1.4a Dimensional Analysis To make conversions between different units we use conversion factors derived from equalities. For example, the equality that relates hours and minutes is 1 hour = 60 minutes. Both 1 hour and 60 minutes represent the same amount of time but with different units. When an equality is written as a ratio, it is called a conversion factor. We use conversion factors to change the units of a quantity without changing the amount using the method called dimensional analysis. In dimensional analysis, the original value is multiplied by the conversion factor. Because both the numerator and the denominator of a conversion factor are equal, the ratio is equal to 1 and multiplying by this ratio does not change the original value; it changes only the units in which it is expressed. For example, you can use the conversion factor 1 hour = 60 minutes to express 48 hours in units of minutes, as shown in the following example

22 EXAMPLE PROBLEM: Use dimensional analysis to convert units. Using dimensional analysis, express 48 hours in units of minutes (1 hour = 60 minutes). SOLUTION: You are asked to express a quantity in different units. You are given a number and an equality. Step 1. Write the quantity in its current units and identify the desired new units: 48 hours conversion factor = minutes Step 2. Identify an equality that relates the current units to the desired units and write it as a ratio. The equality 1 hour = 60 minutes can be written as a ratio in two ways: 1 hour 60 min or 60 min 1 hour Because our goal is to convert units of hours to units of minutes we will use the second form of the conversion factor. Remember that we will multiply the original value by the conversion factor. The second form of the conversion factor will allow the hours units to be cancelled in the multiplication step. Step 3. Multiply the original value by the conversion factor and make sure that the original units are cancelled by the same units in the denominator. The desired units should appear in the numerator. 60 min 48 hours = minutes 1 hour Notice that the amount of time will not change when we perform this multiplication because the conversion factor is derived from an equality and is therefore equal to 1 (60 minutes is equal to 1 hour). Step 4 Perform the multiplication and cancel any units that appear in both the numerator and denominator. Notice that the answer is reported to two significant figures. The equality 1 hour = 60 minutes is exact and therefore has an infinite number of significant figures. The initial value, 48 hours, has two significant figures. Is your answer reasonable? One hour is the equivalent of 60 minutes, so 48 hours is the equivalent of a large number of minutes, much greater than 60 minutes T: Tutorial Assignment Use Dimensional Analysis to Convert Units: Tutorial Assignment 1.4.1: Mastery Assignment Use Dimensional Analysis to Convert Units: Mastery Assignment

23 It is common to combine conversion factors in sequence to perform a unit conversion if you are not provided with a single conversion factor that converts from the given unit to the desired unit, as shown in the following example. EXAMPLE PROBLEM: Use dimensional analysis with more than one conversion factor. Using dimensional analysis, express the volume of soda in a soft drink can (12 oz.) in units of milliliters. (1 gallon = 128 oz.; 1 gallon = L; 1 L = 1000 ml) SOLUTION: You are asked to express a quantity in different units. You are given a number and an equality. Step 1. Write the quantity in its current units and identify the desired new units. 12 oz. conversion factor = ml Step 2. Identify equalities that will convert the given units to the desired units and write them as ratios. First convert units of ounces (oz.) to units of gallons (gal). The next conversion changes units of gallons (gal) to units of liters (L). Finally, convert units of liters (L) to units of milliliters (ml). 1 gal Conversion factor 1: 128 oz L Conversion factor 2: 1 gal 1000 ml Conversion factor 3: 1 L Step 3. Multiply the original value by the conversion factors and make sure that the original units are cancelled by the same units in the denominator. The desired units should appear in the numerator. 1 gal 12 oz L 1000 ml = ml 128 oz. 1 gal 1 L Step 4. Perform the multiplication and division and cancel any units that appear in both the numerator and denominator. Notice that the initial value limits the answer to two significant figures. Is your answer reasonable? One liter is approximately equal to one quart, and a can of soda has a volume that is less than half a quart. Thus, it is reasonable for the volume of a can of soda to be less than 500 ml T: Tutorial Assignment Use Dimensional Analysis with More Than One Conversion Factor: Tutorial Assignment 1.4.2: Mastery Assignment Use Dimensional Analysis with More Than One Conversion Factor: Mastery Assignment

24 1.4b Unit Conversions Using Density Density, the relationship between mass and volume, is also an equality that relates the mass and volume of a substance. Whether working in a chemistry lab or a kitchen (Interactive Figure 1.4.1), you can use the density of a material to convert between units of mass and volume. Interactive Figure Use density in calculations. Some household products with different densities

25 EXAMPLE PROBLEM: Convert units using density. Use the density of diamond (d = 3.52 g/cm3) to determine the volume of a 1.2-carat diamond. (1 carat = g) SOLUTION: You are asked to calculate the volume of a solid from its mass. You are given the mass of the solid and its density. Step 1 Write the quantity in its current units and identify the desired new units. 1.2 carat conversion factor = cm 3 Step 2. Identify equalities that will convert the given units to the desired units and write them as ratios. First convert from carats to units of grams. Next, use density to convert mass units (g) to volume units (cm3) g Conversion factor 1: 1 carat 1 cm 3 Conversion factor 2: 3.52 g Notice that the density conversion factor is inverted. Density can be expressed as an equality (3.54 g = 1 cm 3 ), so it can be written with either mass or volume in the numerator. In this example, the desired unit, cubic centimeters (cm 3 ), appears in the numerator. Step 3. Multiply the original value by the correct conversion factors and make sure that the original unit is cancelled by the same unit in the denominator. The desired unit should appear in the numerator. 1.2 carat g 1 cm3 1 carat 3.52 g = cm3 Step 4. Perform the multiplication and cancel any units that appear in both the numerator and denominator. Is your answer reasonable? While a 1.2-carat diamond is pretty large in the world of diamonds, it does have a relatively small volume. It is reasonable for the volume of this solid to be less than 1 cm T: Tutorial Assignment Convert Units Using Density: Tutorial Assignment 1.4.3: Mastery Assignment Convert Units Using Density: Mastery Assignment

26 Chapter Review Key Concepts 1.1 What Is Chemistry? Chemistry is the study of matter, its transformations, and how it behaves (1.1). Matter is any physical substance that occupies space and has mass (1.1). The macroscopic scale is use to describe matter that can be seen with the naked eye, and the atomic scale is used to describe individual atoms and molecules (1.1a). 1.2 Classification of Matter Matter is described by a collection of characteristics called properties (1.2). An element is a pure substance that cannot be broken down into simpler substances and an atom is the smallest indivisible unit of an element (1.2a). A chemical compound is a substance formed when two or more elements are combined in a defined ratio (1.2a). Molecules are collections of atoms that are held together by chemical bonds (1.2a). A pure substance contains only one type of element or compound and has fixed chemical composition (1.2b). The physical properties of a chemical substance are those that do not change the composition of the material when they are measured, such as physical state (solid, liquid, or gas), melting point, and boiling point (1.2b). The chemical properties of a substance are those that involve a chemical change in the material (1.2b). A solid is a material with a defined shape in which the atoms or molecules are packed together closely. A liquid is a material that flows and takes on the shape of its container because its particles can flow past one another. A gas is a material that has no fixed shape or volume because its particles are widely spaced and move rapidly past one another (1.2b). A change in the physical property of a substance is a physical change, whereas a chemical change involves a change in the chemical composition of the material (1.2b). A mixture is a substance made up of two or more elements or compounds that have not reacted chemically and can be homogeneous (the composition is constant throughout the mixture) or heterogeneous (the composition us not uniform throughout the mixture) (1.2c). 1.3 Units and Measurement SI units are the standard set of scientific units used to report measurements in chemistry (1.3a). Metric prefixes are used with SI units to reflect the relative size of a measured quantity (1.3a). Scientific notation is used to represent very large or very small numbers (1.3b) Temperature is a relative measure of how hot or cold a substance is and is reported using the Fahrenheit, Celsius, or Kelvin temperature scales (1.3c). Absolute zero (0 K) is the lowest temperature possible (1.3c)

27 Density is the ratio of the mass of a substance to its volume (1.3d). Significant figures are the certain and uncertain digits in a measurement (1.3e). When multiplication or division is performed, the number of significant figures in the answer is equal to the number of significant figures in the least certain measurement (1.3f). When addition or subtraction is performed, the number of decimal places in the answer is equal to the number of decimal places in the measurement with the fewest decimal places (1.3f). Precision, how close the values in a set of measurements are to one another, and accuracy, how close a measurement or set of measurements is to a real value, are ways of describing the certainty of a measurement or set of measurements (1.3g). 1.4 Unit Conversions A conversion factor is an equality written in the form of a ratio (1.4a). Dimensional analysis is the use of conversion factors to change the units of a quantity without changing its amount (1.4a). Key Equations Key Terms T (ºC) = T (K) (1.1) 9 T (ºF) = 5 [T (ºC)] + 32 (1.2) mass density = volume (1.3) 1.1 What Is Chemistry? chemistry matter macroscopic scale atomic scale 1.2 Classification of Matter properties element atom chemical compound molecules pure substance physical properties states of matter solid liquid gas melting point freezing point boiling point physical change chemical properties

28 chemical change mixture homogeneous mixture heterogeneous mixture 1.3 Units and Measurement SI units scientific notation length mass temperature absolute zero volume energy density significant figures precision accuracy 1.4 Unit Conversions conversion factor dimensional analysis